Extraction of Metals

Reactivity series (relevant section).

K, Na, Ca, Mg, Al — more reactive than carbon → must use ELECTROLYSIS. C — reference point. Zn, Fe, Sn, Pb, (H), Cu, Ag, Au — less reactive than carbon → can be reduced by carbon.

Why? Reduction by carbon means carbon GIVES UP electrons to the metal oxide. Carbon can do this for metals less reactive than itself (their oxides "want" the electrons more weakly). For more reactive metals, the oxide holds the electrons too tightly — carbon can't reduce them.

Native metals. A few very unreactive metals (Au, Ag) are found in native (uncombined) form in nature — no extraction needed; just physical separation.

Cambridge tip. When asked "why is aluminium extracted by electrolysis?", say: aluminium is more reactive than carbon, so carbon cannot reduce its oxide. Electrolysis is needed.

Inputs (charged from the top of the furnace).

• Haematite ($\text{Fe}_{2}\text{O}_{3}$) — iron ore.
• Coke (carbon, C) — provides heat AND reduces the iron oxide.
• Limestone ($\text{CaCO}_{3}$) — removes acidic impurities.
• Hot air blasted in at the bottom — supplies O₂.

Key reactions inside the furnace.

1. Burning of coke (highly exothermic). $\text{C} + \text{O}_{2} \to \text{CO}_{2}.$

2. Reduction of CO₂ by hot coke (forms CO). $\text{CO}_{2} + \text{C} \to 2\text{CO}.$

3. CO reduces haematite to iron (the main reduction). $\text{Fe}_{2}\text{O}_{3} + 3\text{CO} \to 2\text{Fe} + 3\text{CO}_{2}.$

(At higher in the furnace, carbon also reduces directly: $2\text{Fe}_{2}\text{O}_{3} + 3\text{C} \to 4\text{Fe} + 3\text{CO}_{2}$.)

4. Limestone removes silica (acidic) impurity.

• Limestone decomposes: $\text{CaCO}_{3} \to \text{CaO} + \text{CO}_{2}$.
• CaO reacts with sand (SiO₂): $\text{CaO} + \text{SiO}_{2} \to \text{CaSiO}_{3}$ (slag).

5. Outputs.

• Molten iron (high density) — collected at the bottom, tapped off.
• Slag (calcium silicate) — floats on top of iron, removed separately.
• Waste gases: mostly CO₂ and N₂ (from air), some CO and unreacted gases.

Why limestone? Iron ore contains sandy/silica impurities (acidic). Calcium oxide is BASIC and neutralises them as slag.

Why electrolysis? Aluminium is more reactive than carbon — carbon cannot reduce alumina (Al₂O₃) effectively. Electrolysis pulls electrons in directly.

Why molten cryolite? Pure Al₂O₃ has m.p. $\sim 2050\degree\text{C}$ — far too hot for industry. Dissolving Al₂O₃ in molten cryolite ($\text{Na}_{3}\text{AlF}_{6}$) lowers the working temperature to $\sim 950\degree\text{C}$ — saves enormous amounts of energy.

The cell.

• Steel container lined with carbon (acts as the cathode).
• Carbon anodes dipping into the molten mixture from above.
• Filled with Al₂O₃ dissolved in molten cryolite.

Reactions.

• Cathode ($-$): $\text{Al}^{3+} + 3\text{e}^{-} \to \text{Al}$. Liquid aluminium collects at the bottom.
• Anode ($+$): $2\text{O}^{2-} \to \text{O}_{2} + 4\text{e}^{-}$. Oxygen gas evolves.

Anode wear. At the high temperature, the oxygen REACTS with the carbon anodes: $\text{C} + \text{O}_{2} \to \text{CO}_{2}.$ The anodes BURN AWAY and must be replaced regularly. Adds significant cost.

Energy cost. Aluminium extraction is one of the most energy-intensive industrial processes. That's why ALUMINIUM RECYCLING saves $\sim 95\%$ of the energy compared to extracting fresh metal.

Worked qualitative. Why does the anode in aluminium electrolysis need to be replaced periodically? Anode is carbon. At electrolysis temperature ($\sim 950\degree\text{C}$), the oxygen produced reacts with hot carbon anode → CO₂ → anode burns away.